This short chapter deals with the discovery of two alkali metals, sodium and potassium, and two alkaline-earth metals, magnesium and calcium. They were discovered, directly in a free state, in the first decade of the 19th century.
The compounds of these metals had been known from very remote times and it is hardly possible to establish more or less accurately when common salt, potash, lime, or magnesia came into use. All these compounds had been man's com-panions long before the metals contained in them were discovered.
A. Lavoisier included lime and magnesia into "The Table of Simple Bodies" but excluded potassium and sodium hydroxides believing that they had complex composition and their nature had to be further studied. One might say that history was unjust to these elements, for barium, for instance, was isolated in a metallic state simultaneously with them, but had been discovered much earlier. However, history is a wayward lady. The discovery of sodium, potas-sium, magnepotas-sium, and calcium is interesting in that it was made possible by electric current being successfully used for the first time. This marked the birth of the electrochemi-cal method, a subsidiary to the chemielectrochemi-cal analysis. Subse-quently, electrolysis of melted compounds made it possible to obtain other metals discovered earlier in their compounds.
That is why we considered it justified to devote a separate chapter to the history of sodium, potassium, magnesium, and calcium. The time span in question is two years and H. Davy, one of the founders of electrochemistry, is the main character.
Sodium and Potassium
Man had known sodium and potassium compounds for a very long time. Carbonates of these metals were used in Egypt for laundry. Common salt, one of the most widespread
sodium compounds, was used in foods from time immemorial;
in some countries it was very expensive and sometimes wars were waged for the right to possess salt mines. Sodium carbonate was usually obtained from salt lakes whereas potassium carbonate by leaching plant ash; for this reason the former was named mineral alkali and the latter vegetable alkali. The word "alkali" was introduced by Geber, a me-dieval alchemist, although he made no distinction between the two carbonates. The differences in their nature were first mentioned in 1683. The Dutch scientist I. Bon noted that when soda and potash were used in the similar process, the shapes of the precipitated crystals were different depending on the initial product.
In 1702 G. Stahl noted the difference in crystals of some sodium and potassium compounds. This was an important step in distinguishing between soda and potash. In 1736 the French chemist A. de Monsean proved that soda was always present in common salt, Glauber's salt, and in borax.
Since an acidic constituent of soda was known, the nature of the basic constituent was of great interest. According to Monsean, soda formed Glauber's salt with sulphuric acid, cubic saltpeter (sodium nitrate) with nitric acid, and a variety of sea salt with hydrochloric acid: isn't this reason enough to deduce that soda is the basis of sea salt?
Although chemists had suspected for a long time that al-kali earths were oxides of metals, the nature of soda and potash had not been studied up to the early 19th century.
Even Lavoisier had no definite idea on this subject. He did not know what the basic constituents of soda and potash were and assumed that nitrogen could be a constituent.
This confusion seems to stem from the similarity between the properties of sodium, potassium, and ammonium salts.
Credit for determining these constituents belongs to H. Davy. At first he was dogged by failures: he could not separate metals from soda and potash with the aid of a gal-vanic battery. However, soon the scientist understood his error—he used saturated aqueous solutions but the presence of water hinders decomposition. In October, 1807, Davy decided to melt anhydrous potash, and as soon as he started electrolysis of the alkali hydroxide melt, small balls re-sembling mercury with bright metallic lustre appeared on
H. D A V Y
the negative electrode immersed into the melt. Some of the balls burnt up immediately with an explosion forming bright flame while the others did not burn, but just dimmed and became covered with a white film. Davy concluded that numerous experiments had shown that the balls were the substance which-he had been looking for and this substance was highly inflammable potassium hydroxide.
Davy studied this metal thoroughly and found that when it reacted with water the resulting flame was due to burning of the hydrogen liberated from water. Having studied the metal obtained from potassium hydroxide, H. Davy began to search for sodium hydroxide using the same method and he succeeded in separating another alkali metal. The scientist noted that for its preparation a much more powerful battery
was required than in the experiments with potash. Neverthe-less, the properties of both metals turned out to be similar.
For a short time the scientist carefully studied the prop-erties of potassium and sodium. Some chemists doubted the elemental nature of sodium and potassium believing that they were compounds of alkalis with hydrogen. However, Gay Lussac and Thenard proved convincingly that Datfy had, indeed; obtained simple substances.'
Magnesium
Magnesium compounds such as asbestos, talcum, dolomite, and nephrite have been known from very remote times and used for various purposes. They, however, were not rec-ognized as individual substances but were considered to be varieties of lime.
In 1618 H. Wiker found mineral springs near Epsom in England. In 1695 a salt (magnesium sulphate) with a bitter taste was discovered in the Epsom spring water and later it was used in medicine.
Scientists established that artificial Epsom salt could be prepared by adding sulphuric acid to the mother solution remaining after the purification of salt extracted from sea water. The difference between Epsom and Glauber's (sodium sulphate) salts was established but the difference between lime and white magnesia remained unclear for a long time.
J. Black was the first to establish the different solubilities of these compounds and their sulphates in water. According to G. Newman, magnesium oxide was considered to be white magnesia in contrast to black magnesia, which is pyrolusite.
Metallic magnesium (although not very pure and in a very small amount) was obtained for the first time in 1808 by H. Davy who used the same procedure as that for isolating potassium and sodium. Large amounts of the pure metal were obtained in 1831 by the French chemist A. Bussy. The name of the element is derived from the word "magnesia".
Calcium
Many calcium minerals, for instance, limestone, gypsum, alabaster, that is, mainly, carbonate and sulphate minerals, have been known for a very long time. In the old days people
already knew how to transform limestone into lime by calcination, as was reported by Pliny the Elder. However, it was only in 1755 that J. Black showed that the weight (mass) losses during calcination were completely caused only by the removal of fixed air, i.e. carbon dioxide.
The name "alabaster" served in antiquity to denote two minerals. For one of them (a variety of calcium sulphate) the name survived up to our days, but in Egypt, for example,
"alabaster" meant a variety of calcite (calcium carbonate).
Gypsum has also been used from times immemorial as a construction material. Gypsum-based solutions found application in building pyramids, temples, and other edifices. Theophrastos applied the name "gypsum" to two minerals: gypsum itself and the product of its partial de-hydration. Pure calcium oxide was described by the German chemist I. Pott back in 1746; however, attempts to obtain metal from it with the aid of various reducing agents failed.
The right approach was suggested by H. Davy. First, he attempted to obtain calcium by passing electric current through humid earth insulated from the air by a kerosene layer. (In a similar way he had tried to prepare barium and strontium.) As a result of his experiments, Davy developed the following method of preparing pure alkaline-earth metals. He mixed humid earth with l/3 (by mass) of mercury oxide and placed the mixture into a platinum vessel con-nected to the positive pole of a high-voltage battery. Then he introduced a drop of mercury at the centre of the mixture.
The platinum electrode placed in the drop was connected with the negative pole of the battery. Amalgam obtained in this way was then separated into mercury and silvery-white metal, calcium. Davy prepared pure calcium in 1808.
In the same year J. Berzelius and M. Pontin obtained calcium independently of Davy using a similar method.
The name of the element originates from the Latin word calx, which means "lime".